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Periodic classification of Element
Class XI
Periodic Classification of Element
Chemistry
Question 3.1:
What is the basic theme of organisation in the periodic table?
The basic theme of organisation of elements in the periodic table is to classify the
elements in periods and groups according to their properties. This arrangement makes the
study of elements and their compounds simple and systematic. In the periodic table, elements
with similar properties are placed in the same group.
Question 3.2:
Which important property did Mendeleev use to classify the elements in his periodic table and
did he stick to that?
Mendeleev arranged the elements in his periodic table ordered by atomic weight or mass.
He arranged the elements in periods and groups in order of their increasing atomic
weight. He placed the elements with similar properties in the same group. However, he did not
stick to this arrangement for long. He found out that if the elements were arranged strictly in
order of their increasing atomic weights, then some elements did not fit within this scheme of
classification.
Therefore, he ignored the order of atomic weights in some cases. For example, the atomic
weight of iodine is lower than that of tellurium. Still Mendeleev placed tellurium (in Group
VI) before iodine (in Group VII) simply because iodine’s properties are so similar tofluorine,
chlorine, and bromine.
Question 3.3:
What is the basic difference in approach between the Mendeleev’s Periodic Law and the
Modern Periodic Law?
Mendeleev’s Periodic Law states that the physical and chemical properties of elements are
periodic functions of their atomic weights. On the other hand, the Modern periodic Law states
that the physical and chemical properties of elements are periodic functions of their atomic
numbers.
Question 3.4:
On the basis of quantum numbers, justify that the sixth period of the periodic table should
have 32 elements.
In the periodic table of the elements, a period indicates the value of the principal
quantum number (n) for the outermost shells. Each period begins with the filling of
principal quantum number (n). The value of n for the sixth period is 6.
For n = 6,
azimuthal quantum number (l) can have values of 0, 1, 2, 3, 4.
According to Aufbau’s principle, electrons are added to different orbitals in order of
theirincreasing energies. The energy of the 6d subshell is even higher than that of the 7s
subshell.
In the 6th period, electrons can be filled in only 6s, 4f, 5d, and 6 p subshells. Now, 6s has
one orbital, 4f has seven orbitals, 5d has five orbitals, and 6p has three orbitals. Therefore,
there are a total of sixteen (1 + 7 + 5 + 3 = 16) orbitals available. According to Pauli’s exclusion
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principle, each orbital can accommodate a maximum of 2 electrons. Thus, 16 orbitals can
accommodate a maximum of 32 electrons.Hence, the sixth period of the periodic table should
have 32 elements.
Question 3.5:
In terms of period and group where would you locate the element with Z =114?
Elements with atomic numbers from Z = 87 to Z = 114 are present in the 7th period of the
periodic table. Thus, the element with Z = 114 is present in the 7th period of the periodic
table.
In the 7th period, first two elements with Z = 87 and Z= 88 are s-block elements, the next 14
elements excluding Z = 89 i.e., those with Z = 90 – 103 are f – block elements, ten elements
with Z = 89 and Z = 104 – 112 are d – block elements, and the elements with Z = 113 – 118
are p – block elements. Therefore, the element with Z = 114 is the second p – block element in
the 7th period. Thus, the element with Z = 114 is present in the 7th period and 4th group of the
periodic table.
Question 3.6:
Write the atomic number of the element present in the third period and seventeenth group of
the periodic table.
There are two elements in the 1st period and eight elements in the 2nd period., The third period
starts with the element with Z = 11. Now, there are eight elements in the third period. Thus,
the 3rd period ends with the element with Z = 18 i.e., the element in the 18th group of the
third period has Z = 18. Hence, the element in the 17th group of the third period has atomic
number Z = 17.
Question 3.7:
Which element do you think would have been named by
(i) Lawrence Berkeley Laboratory
(ii) Seaborg’s group? Answer
(i) Lawrencium (Lr) with Z = 103 and Berkelium (Bk) with Z = 97
(ii) Seaborgium (Sg) with Z = 106
Question 3.8:
Why do elements in the same group have similar physical and chemical properties?
The physical and chemical properties of elements depend on the number of valence
electrons. Elements present in the same group have the same number of valence
electrons. Therefore, elements present in the same group have similar physical and
chemical properties.
Question 3.9:
What does atomic radius and ionic radius really mean to you?
Atomic radius is the radius of an atom. It measures the size of an atom. If the element is a metal,
then the atomic radius refers to the metallic radius, and if the element is a non-metal, then it
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refers to the covalent radius. Metallic radius is calculated as half the internuclear distance
separating the metal cores in the metallic crystal. For example, the internuclear distance
between two adjacent copper atoms in solid copper is 256 pm.
Thus, the metallic radius of copper is taken as
Covalent radius is measured as the distance between two atoms when they are found together by
a single bond in a covalent molecule. For example, the distance between two chlorine atoms in
chlorine molecule is 198 pm. Thus, the covalent radius of chlorine istaken as
.
Ionic radius means the radius of an ion (cation or anion). The ionic radii can be calculated by
measuring the distances between the cations and anions in ionic crystals. Since a cation is
formed by removing an electron from an atom, the cation has fewer electrons than the parent
atom resulting in an increase in the effective nuclear charge.
Thus, a cation is smaller than the parent atom. For example, the ionic radius of Na+ion is 95 pm,
whereas the atomic radius of Na atom is 186 pm. On the other hand, an anion is larger in size
than its parent atom. This is because an anion has the same nuclear charge, but more
electrons than the parent atom resulting in an increased repulsion among the electrons and a
decrease in the effective nuclear charge. For example, the ionic radius of F– ion is 136 pm,
whereas the atomic radius of F atom is 64 pm.
Question 3.10:
How does atomic radius vary in a period and in a group? How do you explain the
variation?
Atomic radius generally decreases from left to right across a period. This is because with in
a period, the outer electrons are present in the same valence shell and the atomic number
increases from left to right across a period, resulting in an increased effective nuclear charge.
As a result, the attraction of electrons to the nucleus increases.
On the other hand, the atomic radius generally increases down a group. This is because down a
group, the principal quantum number (n) increases which results in an increase of the distance
between the nucleus and valence electrons.
Question 3.11:
What do you understand by isoelectronic species? Name a species that will be
isoelectronic with each of the following atoms or ions.
(i) F–(ii) Ar (iii) Mg2+ (iv) Rb+
Atoms and ions having the same number of electrons are called isoelectronic species.
(i) F– ion has 9 + 1 = 10 electrons. Thus, the species isoelectronic with it will also have 10
electrons. Some of its isoelectronic species are Na+ ion (11 – 1 = 10 electrons), Ne
(10 electrons), O2– ion (8 + 2 = 10 electrons), and Al3+ ion (13 – 3 = 10 electrons).
(ii) Ar has 18 electrons. Thus, the species isoelectronic with it will also have 18
electrons. Some of its isoelectronic species are S2– ion (16 + 2 = 18 electrons), Cl– ion
(17 + 1 = 18 electrons), K+ ion (19 – 1 = 18 electrons), and Ca2+ ion
(20 – 2 = 18 electrons).
(iii) Mg2+ ion has 12 – 2 = 10 electrons. Thus, the species isoelectronic with it will also have
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10 electrons. Some of its isoelectronic species are F– ion (9 + 1 = 10 electrons), Ne
(10 electrons), O2– ion (8 + 2 = 10 electrons), and Al3+ ion (13 – 3 = 10 electrons).
(iv) Rb+ ion has 37 – 1 = 36 electrons. Thus, the species isoelectronic with it will also have 36
electrons. Some of its isoelectronic species are Br– ion (35 + 1 = 36 electrons), Kr
(36 electrons), and Sr 2+ ion (38 – 2 = 36 electrons).
Question 3.12:
Consider the following species: N3–, O2–, F–, Na+, Mg2+ and Al3+
(a) What is common in them?
(b) Arrange them in the order of increasing ionic radii.
(a) Each of the given species (ions) has the same number of electrons (10 electrons). Hence,
the given species are isoelectronic.
(b) The ionic radii of isoelectronic species increases with a decrease in the magnitudes of nuclear
charge.
The arrangement of the given species in order of their increasing nuclear charge is as
follows:
N3– < O2– < F– < Na+ < Mg2+ < Al3+
Nuclear charge = +7 +8 +9 +11 +12 +13
Therefore, the arrangement of the given species in order of their increasing ionic radii is as
follows:
Al3+ < Mg2+ < Na+ < F– < O2– < N3–
Question 3.13:
Explain why cations are smaller and anions larger in radii than their parent atoms?
A cation has a fewer number of electrons than its parent atom, while its nuclear charge remains
the same. As a result, the attraction of electrons to the nucleus is more in a cation than in its
parent atom. Therefore, a cation is smaller in size than its parent atom. On the other hand, an
anion has one or more electrons than its parent atom, resulting in an increased repulsion among
the electrons and a decrease in the effective nuclear charge. As a result, the distance
between the valence electrons and the nucleus is more in anions than in it’s the parent atom.
Hence, an anion is larger in radius than its parent atom.
Question 3.14:
What is the significance of the terms - ‘isolated gaseous atom’ and ‘ground state’ while
defining the ionization enthalpy and electron gain enthalpy? (Hint: Requirements for comparison
purposes.)
Ionization enthalpy is the energy required to remove an electron from an isolated
gaseous atom in its ground state. Although the atoms are widely separated in the gaseous
state, there are some amounts of attractive forces among the atoms. To determine the
ionization enthalpy, it is impossible to isolate a single atom. But, the force of attraction can be
further reduced by lowering the pressure. For this reason, the term ‘isolated gaseous atom’ is
used in the definition of ionization enthalpy.
Ground state of an atom refers to the most stable state of an atom. If an isolated gaseous
atom is in its ground state, then less amount energy would be required to remove an
electron from it. Therefore, for comparison purposes, ionization enthalpy and electron gain
enthalpy must be determined for an ‘isolated gaseous atom’ and its ‘ground state’.
Question 3.15:
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Energy of an electron in the ground state of the hydrogen atom is –2.18 × 10–18 J.
Calculate the ionization enthalpy of atomic hydrogen in terms of J mol–1.
The energy of an electron in the ground state of the hydrogen atom is –2.18 × 10–18 J. Therefore,
the energy required to remove that electron from the ground state of hydrogen atom is
2.18 × 10–18 J.
Ionization enthalpy of atomic hydrogen = 2.18 × 10–18 J ,Hence, ionization enthalpy of atomic
hydrogen in terms of J mol–1= 2.18 × 10–18 × 6.02 × 1023 J mol–1 = 1.31 × 106 J mol–1
Question 3.16:
Among the second period elements the actual ionization enthalpies are in the order
Li < B < Be < C < O < N < F < Ne.
Explain why
(i) Be has higher ΔiH than B
(ii) O has lower ΔiH than N and F?
(i) During the process of ionization, the electron to be removed from beryllium atom is a 2selectron, whereas the electron to be removed from boron atom is a 2p-electron. Now, 2selectrons are more strongly attached to the nucleus than 2p-electrons. Therefore, more
energy is required to remove a 2s-electron of beryllium than that required to remove a 2pelectron of boron. Hence, beryllium has higher ΔiH than boron.
(ii) In nitrogen, the three 2p-electrons of nitrogen occupy three different atomic orbitals.
However, in oxygen, two of the four 2p-electrons of oxygen occupy the same 2p-orbital. This
results in increased electron-electron repulsion in oxygen atom. As a result, the energy
required to remove the fourth 2p-electron from oxygen is less as compared to the energy
required to remove one of the three 2p-electrons from nitrogen. Hence, oxygen has lower
ΔiH than nitrogen.
Fluorine contains one electron and one proton more than oxygen. As the electron is being
added to the same shell, the increase in nuclear attraction (due to the addition of a proton) is
more than the increase in electronic repulsion (due to the addition of an electron). Therefore,
the valence electrons in fluorine atom experience a more effective nuclear charge than that
experienced by the electrons present in oxygen. As a result, more energy is required to
remove an electron from fluorine atom than that required to remove an electron from oxygen
atom. Hence, oxygen has lower ΔiH than fluorine.
Question 3.17:
How would you explain the fact that the first ionization enthalpy of sodium is lower than that of
magnesium but its second ionization enthalpy is higher than that of magnesium?
The first ionization enthalpy of sodium is more than that of magnesium. This is primarily
because of two reasons:
1. The atomic size of sodium is greater than that of magnesium
2. The effective nuclear charge of magnesium is higher than that of sodium
For these reasons, the energy required to remove an electron from magnesium is more than the
energy required in sodium. Hence, the first ionization enthalpy of sodium is lower than that
of magnesium.
However, the second ionization enthalpy of sodium is higher than that of magnesium. This
is because after losing an electron, sodium attains the stable noble gas
configuration. On the other hand, magnesium, after losing an electron still has one
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electron in the 3s-orbital. In order to attain the stable noble gas configuration, it still has to
lose one more electron. Thus, the energy required to remove the second electron in case of
sodium is much higher than that required in case of magnesium. Hence, the second ionization
enthalpy of sodium is higher than that of magnesium.
Question 3.18:
What are the various factors due to which the ionization enthalpy of the main group elements
tends to decrease down a group?
The factors responsible for the ionization enthalpy of the main group elements to
decrease down a group are listed below:
(i) Increase in the atomic size of elements: As we move down a group, the number of shells
increases. As a result, the atomic size also increases gradually on moving down a group. As the
distance of the valence electrons from the nucleus increases, the electrons are not held very
strongly. Thus, they can be removed easily. Hence, on moving down a group, ionization energy
decreases.
(ii) Increase in the shielding effect: The number of inner shells of electrons increases on moving
down a group. Therefore, the shielding of the valence electrons from the nucleus by the inner
core electrons increases down a group. As a result, the valence electrons are not held very
tightly by the nucleus. Hence, the energy required to remove a valence electron decreases down a
group.
Question 3.19:
The first ionization enthalpy values (in kJmol–1) of group 13 elements are :
B
Al
Ga
In
Tl
801
577
579
558
589
How would you explain this deviation from the general trend?
On moving down a group, ionization enthalpy generally decreases due to an increase in the
atomic size and shielding. Thus, on moving down group 13, ionization enthalpy decreases
from B to Al. But, Ga has higher ionization enthalpy than Al. Al follows mmediately
after s–block elements, whereas Ga follows after d–block elements. The shielding provided by
d-electrons is not very effective. These electrons do not shield the valence electrons very
effectively. As a result, the valence electrons of Ga experience a greater effective nuclear
charge than those of Al. Further, moving from Ga to In, the ionization enthalpy decreases due
to an increase in the atomic size and shielding. But, on moving from In to Tl, the ionization
enthalpy again increases. In the periodic table, Tl follows after 4f and 5d electrons. The
shielding provided by the electrons in both these orbitals is not very effective. Therefore, the
valence electron is held quite strongly by the nucleus. Hence, the ionization energy of Tl is on
the higher side.
Question 3.20:
Which of the following pairs of elements would have a more negative electron gain
enthalpy?
(i) O or F (ii) F or Cl
(i) O and F are present in the same period of the periodic table. An F atom has one proton
and one electron more than O and as an electron is being added to the same shell, the
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atomic size of F is smaller than that of O. As F contains one proton more than O, its nucleus
can attract the incoming electron more strongly in comparison to the nucleus of O atom.
Also, F needs only one more electron to attain the stable noble gas configuration. Hence, the
electron gain enthalpy of F is more negative than that of O.
(ii) F and Cl belong to the same group of the periodic table. The electron gain enthalpy usually
becomes less negative on moving down a group. However, in this case, the value of the electron
gain enthalpy of Cl is more negative than that of F. This is because the atomic size of F is
smaller than that of Cl. In F, the electron will be added to quantum level n = 2, but in Cl, the
electron is added to quantum level n = 3. Therefore, there are less electron- electron repulsions
in Cl and an additional electron can be accommodated easily. Hence, the electron gain enthalpy
of Cl is more negative than that of F.
Question 3.21:
Would you expect the second electron gain enthalpy of O as positive, more negative or less
negative than the first? Justify your answer.
When an electron is added to O atom to form O– ion, energy is released. Thus, the first
electron gain enthalpy of O is negative.
On the other hand, when an electron is added to O– ion to form O2– ion, energy has to be given
out in order to overcome the strong electronic repulsions. Thus, the second electron gain
enthalpy of O is positive.
Question 3.22:
What is the basic
electronegativity?
difference
between
the
terms
electron
gain
enthalpy
and
Electron gain enthalpy is the measure of the tendency of an isolated gaseous atom to accept
an electron, whereas electronegativity is the measure of the tendency of an atom in a chemical
compound to attract a shared pair of electrons.
Question 3.23:
How would you react to the statement that the electronegativity of N on Pauling scale is 3.0 in
all the nitrogen compounds?
Electronegativity of an element is a variable property. It is different in different
compounds. Hence, the statement which says that the electronegativity of N on Pauling scale is
3.0 in all nitrogen compounds is incorrect. The electronegativity of N is different in NH3 and
NO2.
Question 3.24:
Describe the theory associated with the radius of an atom as it
(a) gains an electron
(b) loses an electron
(a) When an atom gains an electron, its size increases. When an electron is added, the number
of electrons goes up by one. This results in an increase in repulsion among the electrons.
However, the number of protons remains the same. As a result, the effective nuclear charge of the
atom decreases and the radius of the atom increases.
(b) When an atom loses an electron, the number of electrons decreases by one while the nuclear
charge remains the same. Therefore, the interelectronic repulsions in the atom decrease. As a
result, the effective nuclear charge increases. Hence, the radius of the atom decreases.
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Question 3.25:
Would you expect the first ionization enthalpies for two isotopes of the same element to be the
same or different? Justify your answer.
The ionization enthalpy of an atom depends on the number of electrons and protons (nuclear
charge) of that atom. Now, the isotopes of an element have the same number of protons and
electrons. Therefore, the first ionization enthalpy for two isotopes of the same element should be
the same.
Question 3.26:
What are the major differences between metals and non-metals?
Metals
1
2
3
4
5
6
7
8
Non–metals
Metals can lose electrons easily.
.
Metals
cannot gain electrons easily.
Metals
generally form ionic compounds.
.
Metals
oxides are basic in nature.
.
Metals
have low ionization enthalpies.
.
Metals
have less negative electron gain
.
Metals
are less electronegative. They are
.
enthalpies.
Metals
have a high reducing
power.
rather
.
electropositive
elements.
.
1
2
3
4
5
6
7
8
Non-metals cannot lose electrons easily.
.
Non-metals
can gain electrons easily.
Non–metals
generallyformcovalent compounds.
.
Non–metallic
oxides are acidic in nature.
.
Non–metals
have high ionizationenthalpies.
.
Non–metals
have high negative electron gain
.
enthalpies.
Non–metals
are electronegative.
.
Non–metals
have a low reducing power.
.
.
Question 3.27:
Use the periodic table to answer the following questions.
(a) Identify an element with five electrons in the outer subshell.
(b) Identify an element that would tend to lose two electrons.
(c) Identify an element that would tend to gain two electrons.
(d) Identify the group having metal, non-metal, liquid as well as gas at the room
temperature.
(a) The electronic configuration of an element having 5 electrons in its outermost subshell
should be ns2 np5. This is the electronic configuration of the halogen group. Thus, the element
can be F, Cl, Br, I, or At.
(b) An element having two valence electrons will lose two electrons easily to attain the stable
noble gas configuration. The general electronic configuration of such an element will be ns2.
This is the electronic configuration of group 2 elements. The elements present in group 2 are
Be, Mg, Ca, Sr, Ba.
(c) An element is likely to gain two electrons if it needs only two electrons to attain the stable
noble gas configuration. Thus, the general electronic configuration of such an element should
be ns2 np4. This is the electronic configuration of the oxygen family.
(d) Group 17 has metal, non–metal, liquid as well as gas at room temperature.
Question 3.28:
The increasing order of reactivity among group 1 elements is Li < Na < K < Rb <Cs
whereas that among group 17 elements is F > CI > Br > I. Explain.
The elements present in group 1 have only 1 valence electron, which they tend to lose. Group 17
elements, on the other hand, need only one electron to attain the noble gas configuration. On
moving down group 1, the ionization enthalpies decrease. This means that the energy required to
lose the valence electron decreases. Thus, reactivity increases on moving down a group. Thus,
the increasing order of reactivity among group 1 elements is as follows:
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Li < Na < K < Rb < Cs
In group 17, as we move down the group from Cl to I, the electron gain enthalpy becomes less
negative i.e., its tendency to gain electrons decreases down group 17. Thus, reactivity decreases
down a group. The electron gain enthalpy of F is less negative than Cl. Still, it is the mostreactive
halogen. This is because of its low bond dissociation energy. Thus, the decreasing order of
reactivity among group 17 elements is as follows:
F > Cl > Br > I
Question 3.29:
Write the general outer electronic configuration of s-, p-, d- and f- block elements.
s block element
p block element
d block element
f block element
ns1-2 , n = 1 to 7
ns2np1-6 ,n= 2 to 7
(n-1) d1- 10 ns0-2 n=4 to 7
(n-2) f1-14(n-1)d0-10 ns2 n= 6 to 7
Question 3.30:
Assign the position of the element having outer electronic configuration
(i) ns2 np4 for n = 3
(ii) (n - 1)d2 ns2 for n = 4, and
(iii) (n - 2) f7 (n - 1)d1 ns2 for n = 6, in the periodic table.
(i) Since n = 3, the element belongs to the 3
rd
period. It is a p–block element since the last
electron occupies the p–orbital.
There are four electrons in the p–orbital. Thus, the corresponding group of the element = Number
of s–block groups + number of d–block groups + number of p–electrons
= 2 + 10 + 4 = 16
Therefore, the element belongs to the 3rd period and 16th group of the periodic table. Hence, the
element is Sulphur.
(ii) Since n = 4, the element belongs to the 4th period. It is a d–block element as d– orbitals are
incompletely filled.There are 2 electrons in the d–orbital. Thus, the corresponding group of the
element = Number of s–block groups + number of d–block groups = 2 + 2 = 4
Therefore, it is a 4th period and 4th group element. Hence, the element is Titanium.
(iii) Since n = 6, the element is present in the 6 period. It is an f –block element as the last
th
electron occupies the f–orbital. It belongs to group 3 of the periodic table since all f-block
elements belong to group 3. Its electronic configuration is [Xe] 4f7 5d1 6s2. Thus, its atomic
number is 54 + 7 + 2 + 1 = 64. Hence, the element is Gadolinium.
Question 3.31:
The first (ΔiH1) and the second (ΔiH2) ionization enthalpies (in kJ mol–1) and the (ΔegH)
electron gain enthalpy (in kJ mol–1) of a few elements are given below:
Elements
I
II
III
IV
V
VI
ΔiH1
520
419
1681
1008
2372
738
ΔiH2
7300
3051
3374
1846
5251
1451
ΔegH
-60
-48
-328
-295
+48
-40
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Which of the above elements is likely to be :
(a) the least reactive element.
(b) the most reactive metal.
(c) the most reactive non-metal.
(d) the least reactive non-metal.
(e) the metal which can form a stable binary halide of the formula MX2, (X=halogen).
(f) the metal which can form a predominantly stable covalent halide of the formula MX
(X=halogen)?
(a) Element V is likely to be the least reactive element. This is because it has the highest first
ionization enthalpy (ΔiH1) and a positive electron gain enthalpy (ΔegH).
(b) Element II is likely to be the most reactive metal as it has the lowest first ionization enthalpy
(ΔiH1) and a low negative electron gain enthalpy (ΔegH).
(c) Element III is likely to be the most reactive non–metal as it has a high first ionization
enthalpy (ΔiH1) and the highest negative electron gain enthalpy (ΔegH).
(d) Element V is likely to be the least reactive non–metal since it has a very high first ionization
enthalpy (ΔiH2) and a positive electron gain enthalpy (ΔegH).
(e) Element VI has a low negative electron gain enthalpy (ΔegH). Thus, it is a metal. Further,
it has the lowest second ionization enthalpy (ΔiH2). Hence, it can form a stable binary halide of
the formula MX2 (X=halogen).
(f) Element I has low first ionization energy and high second ionization energy. Therefore, it
can form a predominantly stable covalent halide of the formula MX (X=halogen).
Question 3.32:
Predict the formula of the stable binary compounds that would be formed by the combination of
the following pairs of elements.
(a) Lithium and oxygen
(c) Aluminium and iodine
(e) Phosphorus and fluorine
(b) Magnesium and nitrogen
(d) Silicon and oxygen
(f) Element 71 and fluorine
(a) LiO2
(b) Mg3N2
(c) AlI3
(d) SiO2
(e) PF3 or PF5
(f) The element with the atomic number 71 is Lutetium (Lu). It has valency 3. Hence, the formula
of the compound is LuF3.
Question 3.33:
In the modern periodic table, the period indicates the value of:
(a) Atomic number
(b) Atomic mass
(c) Principal quantum number
(d) Azimuthal quantum number. Answer
The value of the principal quantum number (n) for the outermost shell or the valence shell
indicates a period in the Modern periodic table.
Question 3.34:
Which of the following statements related to the modern periodic table is incorrect?
(a) The p-block has 6 columns, because a maximum of 6 electrons can occupy all the orbitals in a
p-shell.
(b) The d-block has 8 columns, because a maximum of 8 electrons can occupy all the
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Periodic classification of Element
orbitals in a d-subshell.
(c) Each block contains a number of columns equal to the number of electrons that can occupy
that subshell.
(d) The block indicates value of azimuthal quantum number (l ) for the last subshell that received
electrons in building up the electronic configuration.
The d-block has 10 columns because a maximum of 10 electrons can occupy all the orbitals in a d
subshell.
Question 3.35:
Anything that influences the valence electrons will affect the chemistry of the element.
Which one of the following factors does not affect the valence shell?
(a) Valence principal quantum number (n)
(b) Nuclear charge (Z)
(c) Nuclear mass
(d) Number of core electrons.
Nuclear mass does not affect the valence electrons
Question 3.36:
The size of isoelectronic species — F–, Ne and Na+ is affected by
(a) Nuclear charge (Z )
(b) Valence principal quantum number (n)
(c) Electron-electron interaction in the outer orbitals (d) None of the factors because their size is
the same.
The size of an isoelectronic species increases with a decrease in the nuclear charge (Z). For
example, the order of the increasing nuclear charge of F–, Ne, and Na+ is as follows:
F– < Ne < Na+
Z
9
10
11
Therefore, the order of the increasing size of F–, Ne and Na+ is as follows: Na+ < Ne < F–
Question 3.37:
Which one of the following statements is incorrect in relation to ionization enthalpy?
(a) Ionization enthalpy increases for each successive electron.
(b) The greatest increase in ionization enthalpy is experienced on removal of electron
from core noble gas configuration.
(c) End of valence electrons is marked by a big jump in ionization enthalpy.
(d) Removal of electron from orbitals bearing lower n value is easier than from orbital having
higher n value.
Electrons in orbitals bearing a lower n value are more attracted to the nucleus than electrons in
orbitals bearing a higher n value. Hence, the removal of electrons from orbitals bearing a higher n
value is easier than the removal of electrons from orbitals having a lower n value.
Question 3.38:
Considering the elements B, Al, Mg, and K, the correct order of their metallic character
is:
(a) B > Al > Mg > K
(b) Al > Mg > B > K
(c) Mg > Al > K > B
(d) K > Mg > Al > B
The metallic character of elements decreases from left to right across a period. Thus, the metallic
character of Mg is more than that of Al.
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Periodic classification of Element
The metallic character of elements increases down a group. Thus, the metallic character
of Al is more than that of B.
Considering the above statements, we get K > Mg.
Hence, the correct order of metallic character is K > Mg > Al > B.
Question 3.39:
Considering the elements B, C, N, F, and Si, the correct order of their non-metallic Character is
a B>C>Si>N>F
c. F>N>C>B>Si
b. Si>C>B>N>F
d. F>N>C>Si>B
The non-metallic character of elements increases from left to right across a period. Thus, the
decreasing order of non-metallic character is F > N > C > B. Again, the non-metallic character of
elements decreases down a group. Thus, the decreasing order of non-metallic characters of C and
Si are C > Si. However, Si is less non-metallic than B i.e., B > Si.Hence, the correct order of their
non-metallic characters is F > N > C > B > Si.
Question 3.40:
Considering the elements F, Cl, O and N, the correct order of their chemical reactivity in terms
of oxidizing property is:
a.
c.
F> Cl> O> N
Cl> F> O> N
b.
d.
F> O> Cl> N
O> F> N> Cl
The oxidizing character of elements increases from left to right across a period. Thus, we get the
decreasing order of oxidizing property as F > O > N. Again, the oxidizing character of elements
decreases down a group. Thus, we get F > Cl. However, the oxidizing character of O is more than
that of Cl i.e., O > Cl.Hence, the correct order of chemical reactivity of F, Cl, O, and N in terms
of their oxidizing property is F > O > Cl > N.
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